3.5.4 Transition Metals - Homogeneous catalysis
Here the catalyst is in the same phase as the reactants. In industry it may be the gas phase, for example in the (now outdated) addition polymerisation of ethene by oxygen.
In the laboratory homogeneous catalysis can be demonstrated in solution.
The redox reaction between iodide ions and peroxodisulfate ions is catalysed by iron(II) ions.
Iodide ions are oxidised according to the half-equation:
2I- I2 + 2e
and peroxodisulfate ions are reduced according to the half equation:
S2O82- + 2e 2SO42-
These two half equations can be combined to give an ionic redox equation:
S2O82- + 2I- I2 + 2SO42-
The reaction can be followed by the appearance of iodine using a starch indicator that gets progressively darker blue, or a quantity of thiosulfate ions can be added in the presence of starch indicator and the time taken for the thiosulfate to be consumed by the iodine can be recorded. This is usually called an iodine 'clock' reaction.
I2 + 2S2O32- S4O62- + 2I-
If the number of moles of thiosulfate ions is known, and the time recorded at the instant the blue starch-iodide complex appears, then the average rate of the reaction to that point (in terms of iodine formed) is equal to the half the moles of thiosufate divided by the time in seconds.
Without a catalyst the reaction is slow, but the rate increases considerably in the presence of iron(II) ions. The iron(II) ions are able to change oxidation state to iron(III) assisting in the donation of to the peroxodisulfate ions. The iron(II) ions are then reformed by removing electrons from the iodide ions.
Hence, the iron(II) is able to use its other oxidation state to donate electrons to one species and then to receive electrons from the other reactant.
The reaction between ethandioate ions and manganate(VII) ions
Ethandioate ions have reducing properties with strong oxidising agents such as the manganate(VII) ion in acidic solution. Ethandioate ions are oxidised according to the half equation:
C2O42- 2CO2 + 2e
Manganate(VII) ions are reduced when they behave as an oxidising agent:
MnO4- +8H+ + 5e Mn2+ + 4H2O
These two equations can be combined by multiplying the first equation by 5 and the second equation by 2, before adding them together:
2MnO4- + 16H+ + 5C2O42- 2Mn2+ + 8H2O + 10CO2
The reaction can be followed by the disappearance of the dark purple colour of the maganate(VII) ions (using for example, colorimetry). However, it proceeeds slowly at room temperature at first, but as manganese(II) ions are formed these autocatalyse the reaction and the rate increases.
They are able to do this as manganese has an available oxidation state of 3+. The manganese(II) ions can donate electrons to manganate(VII) and receive them back from ethandioate ions.